Chemistry 9701 (2022, 2023, & 2024): AS Level

Physical chemistry

Inorganic chemistry

Organic chemistry

Analysis

Atomic structure

Atoms, molecules and stoichiometry

Chemical bonding

States of matter 

Chemical energetics

Electrochemistry

Equilibria

Reaction kinetics

The Periodic Table: chemical periodicity

Group 2

Group 17

Nitrogen and sulfur 

An introduction to AS Level organic chemistry

Hydrocarbons

Halogen compounds

Hydroxy compounds 

Carbonyl compounds

Carboxylic acids and derivatives

Nitrogen compounds

Polymerisation

Organic synthesis

Analytical techniques

Particles in the atom and atomic radius

Isotopes

Electrons, energy levels and atomic orbitals: In 1.3 each atom or ion described will be in the ground state. Only the elements hydrogen to krypton will be assessed.

Ionisation energy: In 1.4 each atom or ion described will be in the ground state. Only the elements hydrogen to krypton will be assessed.

Relative masses of atoms and molecules

The mole and the Avogadro constant

Formulae

Reacting masses and volumes (of solutions and gases)

Electronegativity and bonding

Ionic bonding

Metallic bonding

Covalent bonding and coordinate (dative covalent) bonding

Shapes of molecules 

Intermolecular forces, electronegativity and bond properties

Dot-and-cross diagrams

The gaseous state: ideal and real gases and pV = nRT

Bonding and structure

Enthalpy change, ΔH

Hess’s Law

Redox processes: electron transfer and changes in oxidation number (oxidation state)

Chemical equilibria: reversible reactions, dynamic equilibrium 

Brønsted–Lowry theory of acids and bases

Rate of reaction

Effect of temperature on reaction rates and the concept of activation energy

Homogeneous and heterogeneous catalysts

Periodicity of physical properties of the elements in Period 3

Periodicity of chemical properties of the elements in Period 3 

Chemical periodicity of other elements

Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds

Physical properties of the Group 17 elements

The chemical properties of the halogen elements and the hydrogen halides

Some reactions of the halide ions

The reactions of chlorine

Nitrogen and sulfur

Formulae, functional groups and the naming of organic compounds

Characteristic organic reactions

Shapes of organic molecules; σ and π bonds

Isomerism: structural and stereoisomerism

Alkanes

Alkenes

Halogenoalkanes

Alcohols

Aldehydes and ketones

Carboxylic acids

Esters

Primary amines

Nitriles and hydroxynitriles

Addition polymerisation

Organic synthesis

Infrared spectroscopy

Mass spectrometry

understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons; electrons are found in shells in the empty space around the nucleus

identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

understand the terms atomic and proton number; mass and nucleon number 

describe the distribution of mass and charge within an atom

describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field  

determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge

state and explain qualitatively the variations in atomic radius and ionic radius across a period and down a group 

define the term isotope in terms of numbers of protons and neutrons

understand the notation x yA for isotopes, where x is the mass or nucleon number and y is the atomic or proton number  

state that and explain why isotopes of the same element have the same chemical properties

state that and explain why isotopes of the same element have different physical properties, limited to mass and density 

understand the terms:

describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells

describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells

describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital

explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion

determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions: e.g. for Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 (full electronic configuration) or [Ar] 3d6 4s2 (shorthand electronic configuration)

understand and use the electrons in boxes notation

describe and sketch the shapes of s and p orbitals

describe a free radical as a species with one or more unpaired electrons

define and use the term first ionisation energy, IE

construct equations to represent first, second and subsequent ionisation energies

identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table 

identify and explain the variation in successive ionisation energies of an element

understand that ionisation energies are due to the attraction between the nucleus and the outer electron

explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells and spin-pair repulsion

deduce the electronic configurations of elements using successive ionisation energy data

deduce the position of an element in the Periodic Table using successive ionisation energy data

define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom

define relative atomic mass, Ar , relative isotopic mass, relative molecular mass, Mr , and relative formula mass in terms of the unified atomic mass unit

define and use the term mole in terms of the Avogadro constant

write formulae of ionic compounds from ionic charges and oxidation numbers (shown by a Roman numeral), including:

write and construct equations (which should be balanced), including ionic equations (which should not include spectator ions)

define and use the terms empirical and molecular formula

understand and use the terms anhydrous, hydrated and water of crystallisation 

calculate empirical and molecular formulae, using given data 

perform calculations including use of the mole concept, involving:

define electronegativity as the power of an atom to attract electrons to itself 

explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells

state and explain the trends in electronegativity across a period and down a group of the Periodic Table

use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary) 

define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions) 

describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons 

define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons

describe covalent bonds in terms of orbital overlap giving σ and π bonds:

define the terms:  

state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples:

predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples

use the concept of electronegativity to explain bond polarity and dipole moments of molecules

describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces

state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

explain the origin of pressure in a gas in terms of collisions between gas molecules and the wall of the container

understand that ideal gases have zero particle volume and no intermolecular forces of attraction 

state and use the ideal gas equation pV = nRT in calculations, including in the determination of Mr

describe, in simple terms, the lattice structure of a crystalline solid which is:

describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility 

deduce the type of structure and bonding present in a substance from given information

understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)

construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy

define and use the terms:

understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds

use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔHr

understand that some bond energies are exact and some bond energies are averages

calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/n 

apply Hess’s Law to construct simple energy cycles 

carry out calculations using cycles and relevant energy terms, including:

calculate oxidation numbers of elements in compounds and ions 

use changes in oxidation numbers to help balance chemical equations

explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number 

explain and use the terms oxidising agent and reducing agent

use a Roman numeral to indicate the magnitude of the oxidation number of an element

understand what is meant by a reversible reaction

define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change 

use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration, pressure or presence of a catalyst on a system at equilibrium 

deduce expressions for equilibrium constants in terms of concentrations, Kc

use the terms mole fraction and partial pressure

deduce expressions for equilibrium constants in terms of partial pressures, K p (use of the relationship between K p and Kc is not required) 

use the Kc and K p expressions to carry out calculations (such calculations will not require the solving of quadratic equations)

calculate the quantities present at equilibrium, given appropriate data

state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction

describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of dynamic equilibrium in the chemical industry and the application of Le Chatelier’s principle

state the names and formulae of the common acids, limited to hydrochloric acid, HCl, sulfuric acid, H2SO4, nitric acid, HNO3, ethanoic acid, CH3COOH

state the names and formulae of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, KOH, ammonia, NH3

describe the Brønsted–Lowry theory of acids and bases

describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution

appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7

explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity  

understand that neutralisation reactions occur when H+ (aq) and OH– (aq) form H2O(l)

understand that salts are formed in neutralisation reactions

sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis

select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)

explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective collisions

explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction

use experimental data to calculate the rate of a reaction 

define activation energy, EA, as the minimum energy required for a collision to be effective

sketch and use the Boltzmann distribution to explain the significance of activation energy

explain qualitatively, in terms both of the Boltzmann distribution and of frequency of effective collisions, the effect of temperature change on the rate of a reaction

explain and use the terms catalyst and catalysis

describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements 

explain the variation in melting point and electrical conductivity in terms of the structure and bonding of the elements  

describe, and write equations for, the reactions of the elements with oxygen (to give Na2O, MgO, Al 2O3, P4O10, SO2), chlorine (to give NaCl, MgCl 2, Al Cl 3, SiCl 4, PCl 5) and water (Na and Mg only)

state and explain the variation in the oxidation number of the oxides (Na2O, MgO, Al 2O3, P4O10, SO2 and SO3 only) and chlorides (NaCl, MgCl 2, Al Cl 3, SiCl 4, PCl 5 only) in terms of their outer shell (valence shell) electrons

describe, and write equations for, the reactions, if any, of the oxides Na2O, MgO, Al 2O3, SiO2, P4O10, SO2 and SO3 with water including the likely pHs of the solutions obtained 

describe, explain, and write equations for, the acid/base behaviour of the oxides Na2O, MgO, Al 2O3, P4O10, SO2 and SO3and the hydroxides NaOH, Mg(OH)2, Al(OH)3 including, where relevant, amphoteric behaviour in reactions with acids and bases (sodium hydroxide only) 

describe, explain, and write equations for, the reactions of the chlorides NaCl, MgCl 2, Al Cl 3, SiCl 4, PCl 5 with water including the likely pHs of the solutions obtained

explain the variations and trends in 9.2.2, 9.2.3, 9.2.4 and 9.2.5 in terms of bonding and electronegativity

suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties 

predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity

deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

describe, and write equations for, the reactions of the elements with oxygen, water and dilute hydrochloric and sulfuric acids

describe, and write equations for, the reactions of the oxides, hydroxides and carbonates with water and dilute hydrochloric and sulfuric acids

describe, and write equations for, the thermal decomposition of the nitrates and carbonates, to include the trend in thermal stabilities

describe, and make predictions from, the trends in physical and chemical properties of the elements involved in the reactions in 10.1.1 and the compounds involved in 10.1.2, 10.1.3 and 10.1.5

state the variation in the solubilities of the hydroxides and sulfates

describe the colours and the trend in volatility of chlorine, bromine and iodine 

describe and explain the trend in the bond strength of the halogen molecules

interpret the volatility of the elements in terms of instantaneous dipole–induced dipole forces

describe the relative reactivity of the elements as oxidising agents

describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions  

describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths

describe the relative reactivity of halide ions as reducing agents

describe and explain the reactions of halide ions with:

describe and interpret, in terms of changes in oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide and recognise these as disproportionation reactions

explain, including by use of an equation, the use of chlorine in water purification to include the production of the active species HOCl and ClO– which kill bacteria. 

explain the lack of reactivity of nitrogen, with reference to triple bond strength and lack of polarity

describe and explain:

state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines

understand that atmospheric oxides of nitrogen (NO and NO2) can react with unburned hydrocarbons to form peroxyacetyl nitrate, PAN, which is a component of photochemical smog

describe the role of NO and NO2 in the formation of acid rain both directly and in their catalytic role in the oxidation of atmospheric sulfur dioxide 

define the term hydrocarbon as a compound made up of C and H atoms only

understand that alkanes are simple hydrocarbons with no functional group

understand that the compounds in the table on page 26 and 27 contain a functional group which dictates their physical and chemical properties 

interpret and use the general, structural, displayed and skeletal formulae of the classes of compound stated in the table on page 26 and 27

understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups detailed in the table on page 26 and 27, up to six carbon atoms (six plus six for esters, straight chains only for esters and nitriles) 

deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal formula

interpret and use the following terminology associated with types of organic compounds and reactions: 

understand and use the following terminology associated with types of organic mechanisms: 

describe organic molecules as either straight-chained, branched or cyclic

describe and explain the shape of, and bond angles in, molecules containing sp, sp2 and sp3 hybridised atoms

describe the arrangement of σ and π bonds in molecules containing sp, sp2 and sp3 hybridised atoms

understand and use the term planar when describing the arrangement of atoms in organic molecules, for example ethene

describe structural isomerism and its division into chain, positional and functional group isomerism

describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism (use of E, Z nomenclature is acceptable but is not required) 

describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds

explain what is meant by a chiral centre and that such a centre gives rise to two optical isomers (enantiomers) (Candidates should appreciate that compounds can contain more than one chiral centre, but knowledge of meso compounds, or nomenclature such as diastereoisomers is not required)  

identify chiral centres and geometrical (cis/trans) isomerism in a molecule of given structural formula including cyclic compounds

deduce the possible isomers for an organic molecule of known molecular formula 

recall the reactions (reagents and conditions) by which alkanes can be produced:

describe:

describe the mechanism of free-radical substitution with reference to the initiation, propagation and termination steps

suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from heavier crude oil fractions

understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of the C–H bonds and their relative lack of polarity 

recognise the environmental consequences of carbon monoxide, oxides of nitrogen and unburnt hydrocarbons arising from the combustion of alkanes in the internal combustion engine and of their catalytic removal

recall the reactions (including reagents and conditions) by which alkenes can be produced:

describe the following reactions of alkenes: 

describe the use of aqueous bromine to show the presence of a C=C bond 

describe the mechanism of electrophilic addition in alkenes, using bromine / ethene and hydrogen bromide /propene as examples

describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary cations formed during electrophilic addition (this should be used to explain Markovnikov addition) 

recall the reactions (reagents and conditions) by which halogenoalkanes can be produced: 

classify halogenoalkanes into primary, secondary and tertiary 

describe the following nucleophilic substitution reactions:

describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by bromoethane

describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups  

recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure

describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)

recall the reactions (reagents and conditions) by which alcohols can be produced:

describe: 

classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one alcohol group  

deduce the presence of a CH3CH(OH)– group in an alcohol, CH3CH(OH)–R, from its reaction with alkaline I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

explain the acidity of alcohols compared with water

recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced:

describe:

describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones in 17.1.2(b)

describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) to detect the presence of carbonyl compounds

deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (Fehling’s and Tollens’ reagents; ease of oxidation)

deduce the presence of a CH3CO– group in an aldehyde or ketone, CH3CO–R, from its reaction with alkaline I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

recall the reactions by which carboxylic acids can be produced:

describe: 

recall the reaction (reagents and conditions) by which esters can be produced:

describe the hydrolysis of esters by dilute acid and by dilute alkali and heat

recall the reactions by which amines can be produced:

recall the reactions by which nitriles can be produced:

 recall the reactions by which hydroxynitriles can be produced:

describe the hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification to produce a carboxylic acid 

describe addition polymerisation as exemplified by poly(ethene) and poly(chloroethene), PVC

deduce the repeat unit of an addition polymer obtained from a given monomer 

identify the monomer(s) present in a given section of an addition polymer molecule

recognise the difficulty of the disposal of poly(alkene)s, i.e. non-biodegradability and harmful combustion products 

for an organic molecule containing several functional groups:

devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus

analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products

analyse an infrared spectrum of a simple molecule to identify functional groups (see the Data section for the functional groups required)

analyse mass spectra in terms of m/e values and isotopic abundances (knowledge of the working of the mass spectrometer is not required)

calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum

deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum  

suggest the identity of molecules formed by simple fragmentation in a given mass spectrum 

deduce the number of carbon atoms, n, in a compound using the M +1 peak and the formula n = 100 × abundance of M +1 ion 1.1 × abundance of M + ion

deduce the presence of bromine and chlorine atoms in a compound using the M +2 peak

shells, sub-shells and orbitals

principal quantum number (n)

ground state, limited to electronic configuration

the prediction of ionic charge from the position of an element in the Periodic Table

recall of the names and formulae for the following ions: NO3 – , CO3 2–, SO4 2–, OH– , NH4 + , Zn2+, Ag+ , HCO3 – , PO4 3– 

write and construct equations (which should be balanced), including ionic equations (which should not include spectator ions)

use appropriate state symbols in equations  

reacting masses (from formulae and equations) including percentage yield calculations  

volumes of gases (e.g. in the burning of hydrocarbons)

volumes and concentrations of solutions

limiting reagent and excess reagent: (When performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified (see also Mathematical requirements section).)

deduce stoichiometric relationships from calculations such as those in 2.4.1 (a)–(d)

describe covalent bonding in molecules including: hydrogen, H2 • oxygen, O2 • nitrogen, N2 • chlorine, Cl 2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4

understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, SO2, phosphorus pentachloride, PCl 5 , and sulfur hexafluoride, SF6

describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH4 + , and in the Al 2Cl 6 molecule

describe covalent bonds in terms of orbital overlap giving σ and π bonds: σ bonds are formed by direct overlap of orbitals between the bonding atoms • π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond

describe how the σ and π bonds form in molecules including H2, C2H6, C2H4, HCN and N2

use the concept of hybridisation to describe sp, sp2 and sp3 orbitals 

define the terms: • bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state • bond length as the internuclear distance of two covalently bonded atoms

use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

BF3 (trigonal planar, 120°) 

CO2 (linear, 180°)

CH4 (tetrahedral, 109.5°)

NH3 (pyramidal, 107°)

H2O (non-linear, 104.5°)

SF6 (octahedral, 90°)

PF5 (trigonal bipyramidal, 120° and 90°)

describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples

use the concept of hydrogen bonding to explain the anomalous properties of H2O (ice and water):  its relatively high melting and boiling points • its relatively high surface tension • the density of the solid ice compared with the liquid water 

describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces

describe the types of van der Waals’ force: • instantaneous dipole – induced dipole (id-id) force, also called London dispersion forces • permanent dipole – permanent dipole (pd-pd) force, including hydrogen bonding

describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole – permanent dipole force between molecules where hydrogen is bonded to a highly electronegative atom

giant ionic, including sodium chloride and magnesium oxide

simple molecular, including iodine, buckminsterfullerene C60 and ice

giant molecular, including silicon(IV) oxide, graphite and diamond  

giant metallic, including copper

standard conditions (this syllabus assumes that these are 298K and 101 kPa) shown by ⦵.

enthalpy change with particular reference to: reaction, ΔHr , formation, ΔHf , combustion, ΔHc , neutralisation, ΔHneut

determining enthalpy changes that cannot be found by direct experiment

use of bond energy data

understand what is meant by a reversible reaction

understand what is meant by dynamic equilibrium in terms of the rate of forward and reverse reactions being equal and the concentration of reactants and products remaining constant

understand the need for a closed system in order to establish dynamic equilibrium

explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy

explain this catalytic effect in terms of the Boltzmann distribution 

construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an effective catalyst 

aqueous silver ions followed by aqueous ammonia (the formation and formula of the [Ag(NH3) 2] + complex is not required)

concentrated sulfuric acid, to include balanced chemical equations 

the basicity of ammonia, using the Brønsted–Lowry theory

the structure of the ammonium ion and its formation by an acid–base reaction

the displacement of ammonia from ammonium salts by an acid–base reaction 

homologous series  

saturated and unsaturated

homolytic and heterolytic fission

free radical, initiation, propagation, termination (the use of arrows to show movement of single electrons is not required)

nucleophile, electrophile, nucleophilic, electrophilic

addition, substitution, elimination, hydrolysis, condensation

oxidation and reduction: (in equations for organic redox reactions, the symbol [O] can be used to represent one atom of oxygen from an oxidising agent and the symbol [H] one atom of hydrogen from a reducing agent)

free-radical substitution 

electrophilic addition 

nucleophilic substitution

nucleophilic addition

addition of hydrogen to an alkene in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat

cracking of a longer chain alkane, heat with Al 2O3

the complete and incomplete combustion of alkanes

the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane

elimination of HX from a halogenoalkane by ethanolic NaOH and heat

dehydration of an alcohol, by using a heated catalyst (e.g. Al 2O3) or a concentrated acid

cracking of a longer chain alkane

the electrophilic addition of (i) hydrogen in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat (ii) steam, H2O(g) and H3PO4 catalyst (iii) a hydrogen halide, HX(g) at room temperature (iv) a halogen, X2

the oxidation by cold dilute acidified KMnO4 to form the diol

the oxidation by hot concentrated acidified KMnO4 leading to the rupture of the carbon–carbon double bond and the identities of the subsequent products to determine the position of alkene linkages in larger molecules

addition polymerisation exemplified by the reactions of ethene and propene  

the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane

electrophilic addition of an alkene with a halogen, X2, or hydrogen halide, HX(g), at room temperature

substitution of an alcohol, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and heat; or with PCl 5; or with SOCl 2

the reaction with NaOH(aq) and heat to produce an alcohol

the reaction with KCN in ethanol and heat to produce a nitrile

the reaction with NH3 in ethanol heated under pressure to produce an amine

the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as exemplified by bromoethane

electrophilic addition of steam to an alkene, H2O(g) and H3PO4 catalyst

reaction of alkenes with cold dilute acidified potassium manganate(VII) to form a diol

 substitution of a halogenoalkane using NaOH(aq) and heat

reduction of an aldehyde or ketone using NaBH4 or LiAlH4

reduction of a carboxylic acid using LiAlH4

hydrolysis of an ester using dilute acid or dilute alkali and heat

the reaction with oxygen (combustion)

substitution to halogenoalkanes, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and heat; or with PCl 5; or with SOCl 2

the reaction with Na(s) 

oxidation with acidified K2Cr2O7 or acidified KMnO4 to: (i) carbonyl compounds by distillation (ii) carboxylic acids by refluxing (primary alcohols give aldehydes which can be further oxidised to carboxylic acids, secondary alcohols give ketones, tertiary alcohols cannot be oxidised)

dehydration to an alkene, by using a heated catalyst, e.g. Al 2O3 or a concentrated acid

formation of esters by reaction with carboxylic acids and concentrated H2SO4 or H3PO4 as catalyst as exemplified by ethanol

classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one alcohol group  

state characteristic distinguishing reactions, e.g. mild oxidation with acidified K2Cr2O7 , colour change from orange to green

the oxidation of primary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce aldehydes

the oxidation of secondary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce ketones

the reduction of aldehydes and ketones, using NaBH4 or LiAlH4 to produce alcohols

the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce hydroxynitriles exemplified by ethanal and propanone  

oxidation of primary alcohols and aldehydes with acidified K2Cr2O7 or acidified KMnO4 and refluxing

hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification  

hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification

the redox reaction with reactive metals to produce a salt and H2(g)

the neutralisation reaction with alkalis to produce a salt and H2O(l)

the acid–base reaction with carbonates to produce a salt and H2O(l) and CO2(g)

esterification with alcohols with concentrated H2SO4 as catalyst

reduction by LiAlH4 to form a primary alcohol 

the condensation reaction between an alcohol and a carboxylic acid with concentrated H2 SO4 as catalyst

reaction of a halogenoalkane with NH3 in ethanol heated under pressure

reaction of a halogenoalkane with KCN in ethanol and heat

the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat

identify organic functional groups using the reactions in the syllabus

predict properties and reactions